Disinfection by Chemical Methods Ozone and Chlorine Dioxide

To rid drinking water of harmful bacteria and viruses, especially those arising from human and animal fecal matter, by use of a chemical agent requires an oxidizing agent more powerful than 02. In some localities, particularly in France and other parts of western Europe but also in some North American cities—Montreal and Los Angeles are examples—ozone is used for this purpose. Since O3 cannot be stored or shipped because of its very short lifetime, it must be generated on-site by a relatively expensive process involving electrical discharge (20,000 V) in dry air. The resulting ozone-laden air is bubbled through the raw water; about 10 minutes of contact is usually sufficient for disinfection. Since the lifetime of ozone molecules is short, there is no residual protection in the purified water to ensure that it will not be subject to future contamination. Some pollutants in water react with the ozone itself and others with free radicals such as hydroxyl and hydroperoxy (Chapters 1-5) that are produced when ozone reacts with water.

Unfortunately, the reaction of ozone with bromine in water leads to the formation of oxygen-containing organic compounds, particularly those containing the carbonyl group, ^C=0> such as formaldehyde and other low-molecular-weight aldehydes and various other compounds, some of which are toxic. In addition, ozone reacts with bromide ion, Br", present in the water to produce the bromate ion, Br03~, a carcinogen in test animals and probably also in humans. The reaction of ozone with bromide, a natural constituent of water that is often present at ppm concentrations, occurs in several steps; the overall reaction is

The bromate ion produced by ozonation may subsequently react with organic matter in the water to produce toxic organobromine compounds, though experiments have shown that the only brominated product under water treatment conditions is dibromoacetonitrile, CHBr2CN, produced by the reaction of bromate ion with acetonitrile. The MCL (maximum contaminant level) of bromate ion in drinking water is set at 10 ppb (0.010 ppm) by the U.S. EPA. Substances such as bromate ion that are produced during water purification are called disinfection by-products, or DBPs. All known chemical methods of disinfecting water produce DBPs of one type or another.

Similarly, chlorine dioxide gas, C102, is used in more than 300 North American and several thousand European communities to disinfect water. The C102 molecules, themselves free radicals, operate to oxidize organic molecules by extracting electrons from them:

C102 + 4 H+ + 5 e~-» cr + 2 HzO

The organic cations created in the accompanying oxidation half-reaction subsequently react with oxygen and eventually become more fully oxidized. Since chlorine dioxide is not a chlorinating agent—it does not generally introduce chlorine atoms into the substances with which it reacts—and since it oxidizes the dissolved organic matter, much smaller amounts of toxic organic chemical by-products are formed than if molecular chlorine were used (see below).

As is the case with ozone, C102 cannot be stored since it is explosive in the high concentrations that its practical use calls for, so it must be generated on-site. This is accomplished by oxidizing its reduced form, the chlorite ion,

C10i~, from the salt sodium chlorite, NaClO-i: i. ' ' t.

Some of the chlorine dioxide in these processes is converted to chlorate ions, C103". The presence of chlorite and chlorate ions as residuals in the final water has raised health concerns due to their potential toxicity. The U.S. EPA has set an MCL of 1.0 ppm for chlorite ion, and a MRDL (maximum residual disinfectant level) of 0.8 ppm for chlorine dioxide, in drinking water.

Disinfection by Chlorination: History

The most common water purification agent used in North America is hypochlorous acid, HOC1. About half the U.S. population uses surface water, and one-quarter of the population uses groundwater, that is disinfected by HOC1. This neutral, covalent compound kills microorganisms, as it readily passes through their cell membranes. In addition to being effective, disinfection by chlorination is relatively inexpensive. Incorporating a small excess of the chemical in the treated water provides it with residual disinfection power during its subsequent storage and transmission to the consumer. Chlorination is more common than ozonation in North America because generally the raw water is less polluted. Chlorination of public water supplies in the United States, Canada, and Great Britain began in the early years of the twentieth century. For the previous 50 years, chlorination had been practiced on an emergency basis during epidemics caused by water-borne pathogens.

Disinfection by Chlorination: Production of Hypochlorous Acid

Like ozone, HOC1 is not stable in concentrated form and so cannot be stored. For large-scale installations, e.g., municipal water treatment plants, it is generated by dissolving molecular chlorine gas, Cl2, in water. At moderate pH values, the equilibrium in the reaction of chlorine with water lies far to the right and is achieved in a few seconds:

Cl2(g) + H20(aq)^^H0Cl(aq) + H+ + CP

Thus a dilute aqueous solution of chlorine in water contains very little aqueous Cl2 itself. If the pH of the reaction water were allowed to become too high, the result would be the ionization of the weak acid HOC1 to the hypochlorite ion, OC1-, which is less able to penetrate bacteria on account of its electrical charge. Once chlorination is complete, the pH of the water is adjusted upward, if necessary, by the addition of lime, CaO.

In small-scale applications of chlorination, as in swimming pools, the handling of cylinders of Cl2 is inconvenient and dangerous. The chlorine gas can be produced as needed on the spot by the electrolysis of salty water. More commonly, hypochlorous acid instead is generated from the salt calcium hypochlorite, Ca(OCl)2, or' is supplied as an aqueous solution of sodium hypochlorite, NaOCl. In water, an acid-base reaction occurs to convert most of the OOP in these substances to HOC1:

Close control of the pH in an environment like a swimming pool is necessary to avoid the shift to the left side in the position of equilibrium for this reaction that occurs if a very alkaline condition is permitted to prevail. On the other hand, corrosion of pool construction materials can occur in acidic water, so the pH is usually maintained above 7 to prevent such deterioration. Maintenance of an alkaline pH also prevents the conversion of dissolved ammonia, NH3, to the chloramines NH2C1, NHCl2, and especially NC13, which is a powerful eye irritant:

NH3 + 3 HOC1-* NC13 + 3 H20

Significant respiratory and eye irritation problems from exposure to chloramines in the air around indoor swimming pools has been reported when appropriate ventilation is unavailable.

It is desirable to adjust the equilibrium point in the OCP-* HOC1

reaction to favor the predominance of the disinfectant molecular species, HOC1. Since the equilibrium between I lOCl and OCF shifts rapidly in favor of the ion between pH values of 7 and 9, however, the acidity level must be meticulously controlled. Swimming pool acidity can be adjusted by the addition of acid (in the form of sodium bisulfate, NaHS04, which contains the acid HSO4 ) or a base (sodium carbonate, Na2C03) or a buffer (sodium bicarbonate, NaHC03, which contains the amphoteric anion HC03"). Chlorine must be constantly replenished in outdoor pools since UV-B and the short-wavelength components of UV-A light in sunshine are absorbed by and decompose the hypochlorite ion, thereby affecting the equilibrium in the OC1"-* HOC1 process toward the ion:

Continue reading here: Cicr 2 cr o2

Was this article helpful?

0 0